Acid-base%20equilibria%20

1
Acid-base equilibria common ions

• Consider solution containing HF (weak acid) and
  salt NaF
• Species in solution HF, H2O, Na, F-
• What effect does presence of NaF have on
  dissociation equilibrium of HF?
• HF(aq) ? H(aq) F-(aq)

• Use Le Chateliers principle extra F- (from NaF)
  shifts equilibrium to the ..
• Because of this, the H will ..
• What happens to the pH?
• This is the Common Ion Effect

2
Equilibrium calculations involving common ions

• 34.6 g of NH4Cl is added to 3.98 L of a 0.0145 M
  solution of NH3. Kb(NH3) 1.8 x 10-5.
• What is the pH of the original solution before
  the addition of NH4Cl?
• NH3(aq) H2O(l) ? NH4(aq) OH-(aq)
• So whats the pH of the solution after the
  addition of the NH4Cl (assume that the volume
  stays constant)?

3
Effect of adding common ions to weak acids?

• What is the general statement that we can make
  about what happens to the acidity of a weak acid
  upon addition of a common ion?

4
Buffered solutions

• Buffer A solution that resists a change in its
  pH when either H or OH- ions are added
• Very important example blood (cells can only
  survive in very narrow pH range, thus constant
  blood pH is vital.)
• Buffers can contain weak acids salts or weak
  bases salts (solutions can be buffered at
  almost any pH)
• Strong acids / bases cannot be used
• But.just how exactly do buffers work?

5
How does buffering work?

• Buffer contains relatively large concs. of weak
  acid HA and its conjugate base A-.
• Add OH- ions into solution, what happens? (OH- is
  a strong base it will look for H) Weak acid HA
  is best source of H.
• OH- HA ? H2O A-
• OH- ions cannot accumulate (replaced by A- ions)

6
How does buffering work?

• If OH- ions converted to A- ions, how does the pH
  stay so stable? Look at eq. expression for H
  dissociation equilibrium for HA

• Eq. H (and thus pH) determined by ratio HA /
  A-.
• If HA and A- are large relative to OH-
  added, change in HA / A- will be very small.
• Thus, H and pH changes will also be very
  small.
• The essence of buffering HA and A- are very
  large relative to OH- added.

7
The essence of buffering
8
Buffer solution pH calculation

• Buffered solution contains 0.50 M HC2H3O2 (Ka
  1.8x10-5) and 0.50 M NaC2H3O2. Calculate pH of
  solution.
• HC2H3O2(aq) ? C2H3O2-(aq) H(aq)

9
pH changes in buffered solutions

• Calculate change in pH that occurs when 0.010 mol
  solid NaOH added to 1.0 L of buffered solution
  from last question.Compare this with pH change
  that occurs when 0.010 mol solid NaOH added to
  1.0 L of H2O.
• Expected pH of buffered solution will change
  very little pH of water will change by much
  larger amount.

10
How to handle these problems
11
Adding H, rather than OH-

• Exactly the same thinking applied when H added
  to buffered solution of weak acid and conjugate
  base salt. A- (conjugate base) has high affinity
  for added H, and will form weak acid HA H A-
  ? HA
• H ions cannot accumulate (replaced by HA)
• Net change of A- to HA, but if A- and HA are
  very large compared to H, little pH change
  will occur (as expected, in buffered solution).

12
Buffer Capacity

• Amount of H/OH- a buffer can absorb without a
  considerable change in pH
• Consider two 1.0 L ammonia/ammonium buffer
  solutions
• Buffer 1 1.0 M NH3, and 1.0 M NH4
• Ka NH3H / NH4
• H Ka. NH4 / NH3 5.6 x 10-10 M pH ?
• What happens to the pH if 0.1 moles of NaOH is
  added?
• After the reaction with NaOH, calculate the H
  and thus the pH.

13
Buffer Capacity

• Buffer 2 0.01 M NH3, and 0.01 M NH4
• Ka NH3H / NH4
• H Ka. NH4 / NH3 5.6 x 10-10 M pH ?
• What happens to the pH if 0.1 moles of NaOH is
  added?
• After the reaction with NaOH, calculate the H
  and thus the pH.

14
Buffer Capacity

• So - what general statement can be made about
  buffer capacity with respect to the initial
  concentration of the weak acid and its conjugate
  base in the buffer?
• High concentrations
• Low concentrations

15
Henderson-Hasselbalch equation

• Useful equation allows for the calculation of
  buffer pH if the concentration of weak acid (HA)
  and conjugate base (A-)are known.
• A more convenient method of pH calculation will
  give the same answer as with our previous method
  of pH calculation.
• Buffered solution contains 0.50 M HC2H3O2 (Ka
  1.8x10-5) and 0.50 M NaC2H3O2. Calculate pH of
  solution (we did this a few slides ago, but lets
  check that H-H equation gives us the same
  answer).

Henderson-Hasselbalch equation
16
Titrations / pH curves

• Experimental method to determine the
  concentration of an acid (or base).
• Think about this If you were given 25 mL of a
  unknown concentration of HCl, a solution of 0.100
  M NaOH and a burette, how could you determine the
  concentration of the HCl?
• What would happen to the pH of the solution of
  HCl as the NaOH was added to it?

17
Titrations / pH curves
18
Titrations

• How do you know when all the acid has been
  neutralized?
• pH paper
• pH meter
• Or Indicator can be used

• Methyl orange yellow in basic solution and red
  in acidic solution
• Many other indicators available

19
Indicators

• Bromothymol blue (yellow in acidic solution),
  slight green tint as base is added, and blue form
  in basic solution

20
Titration Calculations
21
Solubility Equilibria

• Fluoride used to combat tooth decay
• One product of F- reaction at site of teeth is
  CaF2(s)
• CaF2(s) ? Ca2(aq) 2F-(aq) (dissolving in
  water)
• Consider equlibrium set up between these species
• CaF2(s) ? Ca2(aq) 2F-(aq)
• Equilibrium Expression for this process?
• Ksp Ca2F-2 (Why is CaF2 not included in
  expression?)
• Ksp Solubility Product Constant (solubility
  product)

22
Solubility vs Ksp

• Ksp Equilibrium Constant (solubility product)
• Solubility Equilibrium Position
• Copper (I) bromide has a measured aqueous
  solubility of 2.0 x 10-4 mol/L at 25 C.
  Calculate its Ksp value.
• Equilibrium reaction?
• Equilibrium Expression? Ksp
• (equilibrium concentrations)
• Initial Concentrations?